CHEMISTRY 205 LECTURE
EXAM IV Material - Section 1
Chapter 22
Chromatography refers to any separation method in which the components are distributed between a stationary phase and
a moving (mobile) phase. Substances are separated by differences in their partition (distribution between a stationary phse
and a mobile phase flowing through a column.
I. Principle
The Separation Process: A partitioning between two phases
t0
t1
t2
t3
t4
Mobile Phase
Mixture
(Initial band with a mixture of
three solutes)
Stationary Phase
(Column packing)
To detector/
collector
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I. Classes
A. Adsorption chromatography-The solute is adsorbed onto a solid stationary phase
B. Partition chromatography-A stationary (liquid) phase is coated onto solid particles. The solid particles are called a
solid support. The mobile phase is either a gas or liquid.
C. Ion-exchange chromatography. Ions (either cations or anions) are bonded to a stationary solid phase (usually a
resin). Ionic solute particles of the opposite charge are electrostaticaly attracted to the charged stationary phase.
The mobile phase is a liquid.
D. Molecular exclusion chromatography-This technique separates molecules by size. A mobile(dgas or liquid)
passes through a very porous stationary phase (a gel). Small molecules pass through in the pores of the solid phase
and while large molecules are excluded and quickly stream pass the porous stationary phase.
E. Affinity chromatography-A solid support is coated with a solute molecule which is highly specific for only one type
of molecule.
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III. Gas Chromatography
Soap-bubble
Flow Meter
A. Experimental Set-up
Flow
Splitter
Sample
Injection
Port
Detector
Pressure
Gauge
Bridge
Carrier
Gas
Supply
Pressure
Regulator
Flow
Meter
Recorder
Column
Thermostated
Oven
1. Sample is injected at the injector
A solute or solute mixture (a volatile liquid) is injected with a micro syringe through a rubber septum
2. A hot injector port rapidly volatiles the sample
3. The solute mixture passes through the column
The sample is carried through the column by an inert carrier gas (He, N2, or H2). The column is a stainless steel
or glass tubing packed with a stationary material-all thermostated at a temperature near or above the boiling point
of the sample.
4. Separation
Compounds are separated into bands in the column.
5. Elution
Bands of the separated compounds moved along the column into a detector.
6. Detection
Recording of peaks = chromatogram
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IV. Separation Process
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VI. Calculations
Using the Ratio Method
AREA XStandard
AREA Ystandard
AREA Xunknown
AREA Yunknown
Conc XStandard
Conc Ystandard
=
Conc Xunknown
Conc Yunknown
Problem
1) Ethanol
2) n-Propanol
3) n-Butanol
Peak Area Std.
.373
.407
.425
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Peak Area Unknown
.212
.428
.073
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EXAM IV Material - Section 2
Chapter 19
Spectrophotometry methods uses light to measure the concentrations of certain chemicals
I. Background- Review!!
A. Radiant Energy
1.
2.
3.IR, UV, X-ray, Microwave, and etc.
4. For all radiation:
velocity = c = x
where:
c = 2.998 x 108 m sec-1
frequency (sec-1)




wavelength (m)
where:
E = energy
and
E = h
h = planck's constant
a.
b.
c. Problem: If 500.0 nm, what is theand E?
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B. Electromagnetic Spectrum
Radiant energy is characterized by its wavelength
Wavelength (m)
10-11
Gamma
Rays
1020
10-9
10-5
10-7
X-Rays
Ultra
Violet
1018
1016
Infra Red
10-3
10-1
Microwaves
1014
1012
Frequency (s-1)
101
103
Radio Frequency
1010
108
106
Visible Spectrum
Violet
Indigo
400 nm
Blue
Green
Orange
Yellow
600 nm
500 nm
Red
700 nm
750 nm
1. Absorption by species (molecules, hydrated ions,complexes etc.) in solution.
a. Energy changes
(1) Electronic
(2) Vibrational
(3) Rotational
b. Absorption of light by a Na atom
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104
Absorption cont'd
a) Tuv
b) Tvis
c) TIR
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2. Absorption and color
3. Absorption spectra
C. Spectrophotometers
red
1. Single beam
420
P
Meter
nm
Radiant
Source
Prism
or
Grating
Phototube
Detector
Amplifier
blue
Cell
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2. Double beam
Light
Source
Sample
Cuvet
Scanning
Monochromator
P
Detector
Rotating Mirror
(Beam chopper)
Mirror
Amplifier
Semitransparent
Mirror
Reference
Cuvet
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P0
Mirror
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Recorder
D. Measurements
1. Po, initial radiant power of the light beam
2. P, final radiant power of the light beam
P
3. % T = P
o
x 100
4. Absorbance, A or abs
Po
Psolvent
A = log P  P
solution
E. Beer - Lambert Law
1. %T and Length of Cell
Log
%T
Length
1_
%T
Length
1
a. %T length
b. A  length
c.
Causes:
1. Reflection losses
2. Scattering losses in solution
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F. % T and sample concentration
Log
%T
1_
%T
[X]
[X]
1. A  [X]
Log 1_
%T
%T
%T
1
2. % T  [X]
G. Absorbance is dependant on [X] and length of the cell
H. Beer's law - equation
Po
log P = bc
Where
Po = %T, blank
P = T, sample
 = molar absorptivity in
also:
Po
log P = abc
L cm-1 mol-1
a = absorpitivity used when c is not in moles
b = sample cell length
c = sample concentration
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I. A
100
= log %T
= abc
Aunk
Cx
Problem:
1. A known conc. of Cr(III) has an abs = 0.315
at 570 nm. The [Cr3+] = 4.00 x 10-3 M
2. An unknown concentration of Cr(III) transmits 35.5 % T in the same cuvette
at 570nm. What is the [Cr+3] ?
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J. Limitations to Beer's Law
1.
2.
K. Analysis
1.
2.
3.
4.
L. Wavelength selection for analysis
Key: Use the maximum absorbance for analysis of compounds
1. Spectrum-%T vs absorbance
2. Scan the compound
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M. Analysis of multicomponent systems {Mixtures}
Key: Absorbances at one are additive:
ATotal
= A1 + A2 + A3 + A4 + A5 +....+....An
= 1bc1 + 2bc2 + 3bc4
+....+nbcn
EXPT: A mixture containing "X" and "Y" is analyzed.
Problem: Both absorb at 500 and 700 nm
Part. 1 Determine the absorptivity constants, ax and ay
Assume b = 1 cm
Absorptivity constants for "X"
Absorptivity constants for "Y"
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Part. B Analysis for concentrations for "X" and "Y"
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N. Problems
1. A sample absorb 50% of the monochromatic radiation. How many times longer must the sample path be
made so it absorbs 90 % of the radiation
2. A solution of 5.0 g "Z" (molar mass = 100.0) per L in a 2.0 in cell absorbs 60% of the radiation. Calc the molar
absorbtivity constant.
3. A 1.200 g sample containing 1.00% Mn is oxidized and diluted to 250 mls:
(1) 5 mls is diluted to 100 mls
A = 0.100
(2) 10 mls is diluted to 100 mls
A = 0.180
(3) 20 mls is diluted to 100 mls
A = 0.340
(4) 30 mls is diluted to 100 mls
A = 0.400
An 1.500g unknown sample is oxidized and diluted to 500 mls and a ten ml aliquot is diluted to 100 mls. This
solution gives an A of 0.260.
Calculate the % Mn in the sample.
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4. The equilibrium constant for the reaction: 2 CrO42- + 2 H+ == Cr2O72- + H2O
is 4.2 x 1014. The molar absorptivities at 345 nm for the two principal species in a solution of K2CrO4 are:
= 1.84 x 103 for chromate and  = 10.7 x 102 for dichromate. What is the theoretical absorbance (in a 1.00 cm
cell)for this solution?
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
EXAM IV Material - Section 3
Chapter 14,15, & 16
I. BACKGROUND/REVIEW
Electrochemistry - The interchange of chemical and electrical energy
A. Redox reaction
B. Half reactions
C. Balancing redox equations
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D. Electrical measurements
1. Charge
2. Current
3. Cell potential
a. Definitions
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b. Table of standard potentials (at standard conditions)
Standard. conditions:
Solutes at1 M (solns) { Actually, activity = 1}
Gasses at 1 atm, 760 torr,760 mm
Metals are pure
T° = 25 °C
Table of E° Values
E° (V)
 Na+ + eZn°  Zn+2 + 2e-
2.71
 Cu+2 + 2eFe°  Fe+2 + 2e-
-.34
Na
.76
Cu°
-.77
°E species is measured against:
2H+ + 2e-
 H2 °E = 0
ie. .....all other potentials are measured relative to the H reduction reaction
c. Nerst equations - E for nonstandard conditions
Use the nerst equation if nonstandard conditions exist....ie [species]  1M
E = E° -
A Cc A Dd
0.059
log A aA b
n
A B
E = E° -
[C]c[D]d
0.059
log [A]a[B]b
n
(1) Using the nerst equation
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(2) Calculations:
Calculate the potential for the following
1. Fe2+
 Fe3+
2. Fe2+
 Fe3+
3. Fe2+
 Fe3+
II. Cells
A. Diagram
B. Cell equations:
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C. Danielle cell
John Daniell invented the first battery to generate electricity
Zn(s) + Cu2+
 Zn2+ + Cu(s)
D. Galvanic cells
Spontaneous reactions/ Reactions that go as written:
Oxidation (spontaneous) || Reduction (spontaneous)
D. Electrolytic cells
Electricity is required for the reaction to proceed as written:
Oxidation || Reduction
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E. Galvanic and Electrolytic examples
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F. E calculations
1. Calculate E when the following reaction is half-over chemically:
Cu2+ + Zn°  Zn2+ + Cu°
2. Calculate the electrode potential for platinum electrodes immersed in a solutions that are 0.075 M in Iron (III)
sulfate and 0.060 M in Iron (II) sulfate
3. Calculate the electrode potential for the following cell:
Ag° | AgBr(sat'd) , Br-(0.04M) || H+(1.00 x 10-4) | H2 (0.90 atm), Pt
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G. Equilibrium calculations
1. Calculate the Keq for the following reaction:
Sn2+ + Cr2+

Sn° + Cr3+
2. Calculate the Ksp of Ag2C2O4 if the cell has a potential of + 0.277
3. The Ksp fosr ZnCO3 is 1.26 x 10-10 Calculate the E° for the following process:
ZnCO3
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 Zn° + CO32-
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Chapter 16
I. PROBLEMS:
Key:
1.
Balance the redox equation first! - Use Ap34
The Sb(III) in a 1.080 g sample of stibnite ore required a 41.67 ml titration with
0.03134 M I2
2.
.
Calculate the percentage stibnite as Sb2O3 in the sample
A 16.00 g sample containing the mineral tellurite was dissolved and then treated
with 75.00 ml of 0.02000 M K2Cr2O7.
3 TeO2
+
Cr2O72-
+
8 H+

3 H2TeO4
+
2 Cr3+
+
H 2O
Upon completion of the reaction, the excess dichromate required a 50.00 ml back-titration
with 0.0.01500M Fe2+.
Calculate the percentage of TeO2 in the sample.
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II. A REDOX TITRATION CURVE
Problem 1: Derive a titration curve for the titration of 30.00 ml of 0.100 M Fe2+ titrated with 0.1000M Ce4+
Construction of Titration Curve
Calculation of the volume of base needed to reach the equivalence point
Region 1 = BEFORE THE EQUIVALENCE POINT
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Region 2 = THE EQUIVALENCE POINT
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Region 3 = AFTER THE EQUIVALENCE POINT
Graph
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III. REDOX INDICATORS
Redox indicator changes color from one oxidation state to another
Ex. Ferroin
(Phen)3Fe3+ + e-

(Phen)3Fe2+
A. Theory
E = E° -
[Inred]
0.05916
log
n
[Inox]
Typically, a color change from the oxidized form of the indicator to the reduced form requires a
in the concentration ratio from:
[Inred]
1
[Inox] < 10
to
[Inred]
[Inox]
10
< 1
B. Types of indicators
1. Auto indicators:
a.
b.
c.
2. Specific Indicators:
a.
b.
C. Titrants
1. Cerium
2. Potassium permanganate
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change
3. Sodium thiosulfate
4. Iron
5. Potassium dichromate
D. IODINE TITRATIONS
1. Solubility of I2
2. Reduction reaction
3. Direct titrations
4. Starch indicator
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